Bonding

1. Bonding

2. Metallic Bonding • In a liquid or solid state, metals readily give up electrons • When only other metal atoms are around, electrons are not accepted and held, they are free to move • Free moving electrons called an electron sea—reason for metallic properties of luster, malleability, ductility, and conductivity

3. Determine the electronegativity difference, bond type, and more electronegative element with respect to the following atoms. • H and F • Br and Br • Al and S • Na and S • At and Cl • Ba and O • Si and O • Cr and Cl • Fe and I • K and Br

4. Review: bond types, bond energy • There are 2 types of bonding: ionic, covalent (metallic not considered) • ionic = stealing of electrons to form + and – ions. + and – then attract • Covalent = sharing of electrons • Remember there is no clear dividing line. What causes atoms to form molecules? • Basically, all things that happen spontaneously are energetically favorable • Something must be energetically favorable about atoms coming together as molecules (ie: by bonding atoms can lower their energy status

5. An ionic bond is the chemical bond that results from the electrostatic attraction between positive ions (cations) and negative ions (anions) and a lattice structure results in which the cations are surrounded by anions and the anions are surrounded by cations. An ionic bond is the chemical bond that results from the electrostatic attraction between positive ions (cations) and negative ions (anions) and a lattice structure results in which the cations are surrounded by anions and the anions are surrounded by cations.

6. Na+ Na+ Na + e– Na+ Cl– Cl– Cl Cl– + e– + Ionic bonding Ionic bonding involves 3 steps (3 energies) 1) loss of an electron(s) by one element, 2) gain of electron(s) by a second element, 3) attraction between positive and negative Ionization energy Electron affinity Lattice energy

7. Ionic bonding: energies • By convention, a requirement for energy is given a + sign (we have to put energy in) and is called endothermic, a release of energy is given a – sign and is called exothermic.

8. Na+ Na+ Na + e– Na+ Cl– Cl– Cl Cl– + e– + Ionic bonding Ionic bonding involves 3 steps (3 energies) 1) loss of an electron(s) by one element, 2) gain of electron(s) by a second element, 3) attraction between positive and negative Ionization energy + 496 Electron affinity – 349 Lattice energy – 766

9. Ionic bonding: energies • By convention, a requirement for energy is given a + sign (we have to put energy in) and is called endothermic, a release of energy is given a – sign and is called exothermic. • Problem: the sum is +147. A spontaneous change must involve a net lowering of energy • Solution: the lattice energy provides the energy needed • Note that although we represent this as a three step process it actual occurs all at once

10. Review Questions • A metal + non-metal gives what kind of bond? • What is necessary for any stable compound to form from its elements? • Define cation. Define anion. • List the 3 energies involved in forming an ionic bond. • What term describes a release of energy? • Is breaking a bond endothermic or exothermic? • Define lattice energy. • Explain why metals form cations and non-metals form anions. • Explain why calcium exists as Ca2+ but not as Ca3+ in ionic compounds. • Explain why most transition metals form a 2+ ion.

11. Answers • Ionic • There must be a net lowering of energy • Cation: a positively charged ion, Anion: a negatively charged ion • Ionization energy, electron affinity, lattice energy • Exothermic • Endothermic • The energy released by the imaginary process in which isolated ions come together to form a crystal of an ionic compound

12. Metals form cations because they have small IEs (and EA), non-metals form anions because they have large EAs (and IE). These trends are energetically favorable. • The first two electrons from the 4s subshell are easily lost (they can be made up for by the lattice energy). Losing a third electron is not energetically favorable because of the large third IE for Ca (the energy required to remove the third electron can not be made up by the lattice energy. • The loss of two electrons from the s subshell accounts for the typical 2+ charge of the transition elements.

13. Covalent bonding • Just as with ionic bonds, covalent bonds must involve a net lowering of energy • We can explain this net lowering of energy in two ways: 1) visualizing the combination of attraction as two atoms approach each other 2) drawing and combining orbital diagrams

14. + – – – – + As atoms approach • Recall that EA for all atoms, except the noble gasses are negative • In other words we have no trouble adding electrons to atoms • The attraction for electrons is not limited to free electrons, but also involves electrons that are part of other atoms. • Thus, atoms are pulled toward each other • How far they are pulled together will depend on a balance of attraction (nucleus to electrons) and repulsion (nucleus to nucleus and electrons to electrons)

15. There are two types of covalent bonds. In a non-polar covalent bond the electrons are shared equally between the two atoms. (The nuclei and core electrons are indicated by the blue spheres and the bonding electrons are indicated by the lavender dots.) • The electrons move around the nuclei with the electrons generating temporary positive and negative charges within the molecule.

16. In a polar covalent bond the electrons are shared unequally between the two atoms. In this situation, one atom has a greater ability to pull the bonding electrons towards it and is said to be more electronegative. (The green sphere represents the more electronegative element.) • Again, the electrons move around the nuclei with the electrons spending the majority of the time near the more electronegative element. This generates a partial negative charge near the more electronegative element and a partial positive charge near the less electronegative element.

17. Molecular polarity is dependent on bond polarity and the molecular geometry. For small molecules: If all the regions surrounding an atom are similar in their electronegativities, the molecule will be non-polar. If the regions are different, then the molecule will be polar. • For large molecules there may be polar and non-polar regions.

18. Lewis Structures • The first step in drawing Lewis structures is to determine the number of electrons to be used to connect the atoms. This is done by simply adding up the number of valence electrons of the atoms in the molecule. Consider carbon dioxide CO2 • carbon (C) has four valence electrons x 1 carbon = 4 e-oxygen (O) has six valence electrons x 2 oxygens = 12 e-There are a total of 16 e- to be placed in the Lewis structure.

19. 2. Connect the central atom to the other atoms in the molecule with single bonds. Carbon is the central atom, the two oxygens are bound to it and electrons are added to fulfill the octets of the outer atoms. (: or means shared pair of electrons) Some atoms can form multiple bonds (i.e.: share more than one pair of electrons )

20. 3. Complete the valence shell of the outer atoms in the molecule.

21. 4. Place any remaining electrons on the central atom. • there are no more available electrons • If the valence shell of the central atom is not complete, use a lone pair on one of the outer atoms to form a double bond between that outer atom and the central atom. Continue this process of making multiple bonds between the outer atoms and the central atom until the valence shell of the central atom is complete.

22. becomes The central atom is still electron deficient, so share another pair. becomes

23. 5. Double check to make sure that you have used the correct number of electrons in the Lewis structure and that no atom that cannot exceed its valence shell, does not. The best Lewis structure that can be drawn for carbon dioxide is:

24. What Is VSEPR? • The Valence Shell Electron Pair Repulsion (VSEPR) model: • is based on the number of regions of high electron density around a central atom. • can be used to predict structures of molecules or ions that contain only non-metals by minimizing the electrostatic repulsion between the regions of high electron density. • can also be used to predict structures of molecules or ions that contain multiple bonds or unpaired electrons. • does fail in some cases.

25. Intermolecular forces • Dipole—dipole • A dipole is created by equal but opposite charges that are separated by a short distance • The direction of a dipole is from its positive pole to its negative pole used to represent • Attraction between dipoles of polar molecules called dipole-dipole force

26. Polar molecules can induce a temporary attraction in nonpolar molecules • Example: • > >

27. Hydrogen bonding • The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule • Represented by dotted lines connecting the bonded hydrogen to the unshared pair of the electronegative atom to which it is attracted

29. London dispersion forces • The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles • Forces between all atoms and molecules (including Noble gases) • Increase with atomic number

30. Boiling Points • The strength of the intermolecular forces is directly correlated to the boiling points. • The higher the bp, the more energy required to break the intermolecular forces • Higher bp, stronger the intermolecular forces